1. The Periodic Table

2. Periodic Table of Elements
There are 117 elements (January, 2007)
Your table contains 113
94 of the elements are naturally occurring, the rest are man-made
Most of the elements were discovered between

3. History/Development
Development of the table has occurred over 300 years and continues today
Dmitri Mendeleev is commonly credited with creating the periodic table in 1869

4. Original Table (Mendeleev)
Classified elements in horizontal rows based on their atomic mass
When there was a repeat of properties the elements were placed in the next row
Concluded that similar properties appear at regular intervals when the elements are listed in order of increasing atomic mass

5. Original Table (Mendeleev)
Elements with similar properties were located in the same vertical columns, if no known element had the expected properties to fit the particular space he left that space empty
Assumed that elements not yet discovered would fit into the empty spaces
Predicted some properties of these unknown elements (gallium, scandium, and germanium)

7. Modern Periodic Law
Henry Moselely bombarded elements with high speed electrons and they emitted X-rays with a certain wavelength
He found that each element differed by one proton
Concluded that the regularity, or periodicity, of the properties is a function of the atomic number (Modern Periodic Law)

8. Questions
The modern periodic table is arranged by increasing _______________________.
The original periodic table was arranged by increasing _____________________.

9. Organization of the Modern Periodic Table
The periodic table is arranged in order of increasing ATOMIC NUMBER
Horizontal rows are called PERIODS
There are 7 periods
Vertical columns are called GROUPS
There are 18 groups

10. Periods Horizontal Rows
The number of each period indicates the principle energy level in which the valence electrons are located
The number of valence electrons increases as you go from left to right
The properties of the elements change systematically through a period

11. Groups (Families)
Vertical Columns
The outermost shell of an atom contains the same number and arrangement of valence electrons
Elements that have similar chemical properties are located in the same group

12. Group/Period Examples
Are elements in the same period or group more similar? Explain why.
Which elements have the most similar chemical properties?
K and Na
K and Cl
K and Ca
K and S
3. Why do elements of a given group on the periodic table show similar chemical properties?

13. States of Matter Solid Liquid Gas
The majority of the elements are solids
Liquid
The only liquids are Hg and Br, which are found on the right
Gas
H, O, N, F, Cl and the Noble Gases (Group 18), located on the right

14. Nonmetals
Metalloids
Metals

15. Metals/Nonmetals/Metalloids
Where are the metals located?
Where are the nonmetals located?
Where are the metalloids located?
List all of the Metalloids

16. Properties of Metals and Nonmetals
Lose electrons to form positive ions
Solids at STP
Except Hg
High melting and boiling points
Good thermal (heat) and electrical conductors
Luster (shine)
Malleable (bendable)
Ductile – can be made into wires
Gain electrons to form negative ions
Solids, 1 liquid (Br), gases at STP
Low melting and boiling points
Poor conductors of heat and electricity
Dull
Brittle
Not ductile

17. Metalloids Properties of both metals and nonmetals
B, Si, As, Te, Ge, and Sb (front + middle back of staircase)

18. Metal/Nonmetal Questions
1. Atoms of metals tend to
a. lose electrons to form positive ions
b. lose electrons to form negative ions
c. gain electrons to form positive ions
d. gain electrons to form negative ions
2. The majority of elements on the table are classified as
a. metals
b. nonmetals
c. metalloids
3. Which property is generally characteristic of metallic elements?
a. low electrical conductivity
b. high heat conductivity
c. existence as brittle solids
d. low melting points

19. Metal/Nonmetal Examples
4. At room temperature, which substance is the best conductor of electricity?
a. nitrogen
b. neon
c. sulfur
d. silver
5. Which element is brittle in the solid phase and a poor conductor of electricity?
a. calcium
b. strontium
d. copper
6. The majority of elements on the table are in what physical state at STP?
a. solid
b. liquid
c. gas

20. Atomic Radius ½ the distance between any two nuclei
Given on Reference Table S in picometers
1pm = 1x10-12m

21. Trend within a Period (Left to Right)
Atomic Radius decreases
As you move across a period the number of protons increases, resulting in a stronger nuclear charge therefore, electrons are pulled closer to the nucleus

22. Trend within a Group (Top to Bottom)
Atomic Radius increases
Remember period number = PEL
As you move down a group there are additional rings, therefore the valence electrons are further away from the nucleus, resulting in a larger radius

23. Atomic Radius Examples
Which sequence of elements is arranged in order of decreasing atomic radii?
Al, Si, P
Li, Na, K
Cl, Br, I
N, C, B
2. What is the radius of Ca?
3. What is the radius of Sr?
4. Explain why Sr has a larger atomic radius than Ca.

24. Ionic Radius (IR)
Radius that results from the loss or gain of electrons

25. Metals (Left Side)
Tend to lose 1 or more electrons when forming positive ions
Radius will decrease
Ionic Radius < Atomic Radius
Ex: Na+ is smaller than Na

26. Non-metals (Right Side)
Tend to gain 1 or more electrons when forming negative ions
Radius will increase
Ionic Radius > Atomic Radius
Ex: Cl- is larger than Cl

27. Ionic Radius Examples Which ion has the largest radius?
a. Na+ b. Mg2+ c. K+ d. Ca2+
2. Which of the following elements has an ionic radius smaller than its atomic radius?
Neon b. Nitrogen c. Sodium d. Sulfur
3. The Na+ ion has a smaller radius than the Ne atom, even though they both contain 10 electrons. Explain why this is so.

28. Ionization Energy
The energy required to remove the most loosely bound electron from an atom
Low IE = greater tendency to lose electrons and form positive ions
High IE = greater tendency to gain electrons and form negative ions
Given on Reference Table S in kilojoules per mole (kJ/mol)

29. IE within a Period (Left to Right)
IE increases
Number of protons increases, resulting in a stronger nuclear charge
The nucleus has a better hold on the electrons, therefore more energy is required to remove an electron

30. IE within a Group (Top to Bottom)
IE decreases
The principle energy levels increase, so the valence electrons are further away
Protons cannot hold onto the valence electrons as well, therefore less energy is required to remove an electron

31. Electronegativity (e-neg)
Scale that measures the ability of an atom to attract electrons from another atom
Reference Table S
Scale ranges from
Fluorine is the highest = 4.0
Difference between electronegativity between two atoms can be used to determine the type of bond

32. Trend within a Period (Left to Right)
Electronegativity Increases
More protons, resulting in a stronger nucleus, therefore the nucleus is better able to attract electrons

33. Trend within a Group (Top to Bottom)
Electronegativity decreases
The atom is larger, so the nucleus is further away from the valence shell, therefore the nucleus is less able to attract electrons

34. IE/e-neg Examples Which element will lose electrons the easiest?
a. Na b. Cl c. K d. F
2. Which element would be most likely to gain electrons?
a. Na b.Cl c. K d. F

35. Reactivity Metals Nonmetals
Most reactive = loses electrons the easiest (low ionization energy)
Lower left corner of the table (Fr)
Nonmetals
Most reactive = gains electrons the easiest (high electronegativity)
Upper right corner of the table (F, Cl, O) – not group 18

36. Periodic Trends Questions
What is the ionization energy of K?
What is the ionization energy of Ca?
Explain why K has a lower ionization energy than Ca.
According to the reference table, which of the following elements has the smallest radius?
a. Ni b. Co c. Ca d. K

37. Periodic Trends Questions
5. An element with high ionization energy would most likely be?
a. A nonmetal with low electronegativity
b. A nonmetal with high electronegativity
c. A metal with low electronegativity
d. A metal with high electronegativity
6. What happens to S when it becomes S2-?
a. It loses two electrons and the radius increases
b. It loses two electrons and the radius decreases
c. It gains two electrons and the radius increases
d. It gains two electrons and the radius decreases

38. Periodic Trends Questions
7. Which of the following would have the largest radius?
a. Na b. Na c. Cl Cl-1
8. Which of the following has the greater ionization energy, Na or Na+? Explain your answer.

39. Groups 1 and 2 Properties: typical metallic characteristics
High reactivity (valence electrons are easily lost)
Only occur in nature as compounds
Reactivity increases as you move down the group
Group 1 elements are more reactive than group 2 elements
For metals – Low IE = high reactivity (electrons are easily lost) – Fr is the most reactive
Na and Water
K and Water

40. Examples: Which atom is the most reactive?
Na b. Mg c. K d. Ca
Which group 15 element has the least metallic character?
N b. P c. As d. Sb
3. Explain why reactivity increases as you move down group 1.

41. Group 17 - Halogens Typical Nonmetals
High electronegativity – F is the highest
High ionization energy
Are so reactive that they cannot exist as in the monoatomic form
Exist in nature as diatomics (HOFBrINCl)
F2 and Cl2 are gases
Br2 is a liquid
I2 and At2 are solids
F is the most reactive nonmetal
For nonmetals – high electronegativity = high reactivity

42. Group 18 – Noble Gases Exist as gases at STP
Exist as monatomic molecules (not combined with anything
Example: He, Ne, Ar
The outermost ring is complete, therefore they are VERY STABLE (unreactive)

43. Groups 3-12 (Transition Elements)
Hard solids
High melting points (except mercury)
Multiple Positive Oxidation States
They can react with electrons from both s and d sublevels
Different numbers of electrons can be lost
Colored Ions
Easily excited (since d and s sublevels are close)

44. Group Examples
The presence of which ion usually produces a colored solution?
a. K+ b. F- c. Fe2+ d. S2-
Which solution would be colored?
a. KNO3 b. Ca(NO3)2 c. Cu(NO3)2 d. Al(NO3)3
Which element at STP exists as monatomic molecules?
a. Ne b. N c. Cl d. O
Which is a solid at STP?
a. F2 b. Cl2 c. Br2 d. I2
Which element in Period 4 is classified as an active nonmetal?
a. Ga b. Ge c. Br Kr

45. Group Examples a. Bi b. As c. P d. N
Which noble gas would most likely form a compound with fluorine?
a. He b. Ne c. Ar d. Kr
Which element in Period 3 is the most reactive metal?
a. Na b. Mg c. N d. Cl
Which element in Group 15 has the most metallic character?
a. Bi b. As c. P d. N

46. Group Examples
Why is hydrogen not considered to be a member of Group 1?
Why is hydrogen considered to be a member of Group 1?
Why is it unlikely for sodium to form the Na2+ ion?

47. Allotropes
Some nonmetals can exist in 2 or more forms in the same phase
Allotropes have different physical and chemical properties because their atoms are arranged differently
Examples:
oxygen and ozone (O2 and O3)
Graphite and Diamond (carbon)
Diamond
Graphite